table of contents - papacambridge · 2020. 5. 2. · covalent compounds are made of molecules which...

TABLE OF CONTENTS 3

CHAPTER 1

Atoms, Molecules & Stoichiometry

3 CHAPTER 2

Atomic Structure

5 CHAPTER 3

Chemical Bonding

7 CHAPTER 4

States of Matter

9 CHAPTER 5

Chemical Energetics

10 CHAPTER 6

Electrochemistry

11 CHAPTER 7

Equilibria

12 CHAPTER 8

Reaction Kinetics

14 CHAPTER 9

Chemical Periodicity

16 CHAPTER 10

Group II – Alkaline Earth Metals

CIE AS-LEVEL CHEMISTRY//9701

PAGE 2 OF 30

17 CHAPTER 11

Group 17 – Halogens

18 CHAPTER 12

Nitrogen & Sulphur

18 CHAPTER 13

Introduction to Organic Chemistry

22 CHAPTER 14

Hydrocarbons

25 CHAPTER 15

Halogen Derivatives

26 CHAPTER 16

Hydroxy Compounds

28 CHAPTER 17

Carbonyl Compounds

29 CHAPTER 18

Carboxylic Acids & Derivatives

29 CHAPTER 19

Analytical Techniques


PAGE 3 OF 30

1. ATOMS, MOLECULES AND STOICHIOMETRY

1.1 Relative Mass

Re

lati

ve

Atomic

mass (Ar):

weighted average mass

of an atom compared

with 12C

where one

atom of 12C

has mass of

exactly 12

units

Molecular

mass (Mr):

mass of a molecule

Formula

mass:

mass of one formula

unit of a compound

Isotopic

mass:

mass of a particular

isotope of an element

1.2 The Mole

Mole: amount of substance that has the same number

of particles (atoms, ions, molecules or electrons) as

there are atoms in exactly 12g of the carbon-12 isotope.

Avogadro’s constant: number of atoms, ions, molecules

or electrons in a mole = 6.02 10

1.3 Mass Spectra

Abundance of isotopes can be represented on a

mass spectra diagram

ℎ ℎ 100%

100

1.4 Empirical and Molecular Formulae

Empirical formula: gives simplest ratio of different

atoms present in a molecule

Molecular formula: gives actual numbers of each type of

atom in a molecule

Molecular formula can be calculated using the Mr of a

compound and its empirical formula

Where

% . 100%

1.5 Calculations involving Mole Concept

24

Formula applies to gases at r.t.p.

Unit of volume is and 1000 1

Concentration unit

2. ATOMIC STRUCTURE

2.1 Subatomic Particle

Subatomic

Particle

Relative

Charge

Relative mass/

a.m.u

Protons (P) +1 1

Neutrons (n) 0 1

Electrons (e-) -1 1/1840

2.2 Behavior of a Beam of Subatomic

Particles

Protons: positively charged ∴ deflected to -ve pole Neutrons: no charge ∴ not deflected Electrons: negatively charged ∴ deflected to +ve pole e- lighter than P+ ∴ deflected at greater angle 2.3 Protons, neutrons and electrons

Mass concentrated within centre; nucleus

An atom is electrically neutral; P+ = e-

Atomic no. or proton no. (Z) = no. of protons

Atomic mass or nucleon no. (A) = no. of P + N

Isoelectronic ions: ions having same no. of e-s

Isotopes: are atoms of the same element with the same

proton number but different number of neutrons

2.4 Electronic Configuration

Electrons are arranged in energy levels called shells

Each shell is described by a principle quantum no. (P.Q)

As the P.Q. increases, energy of shell increases

Inside the shell there are subshells: , , and Orbital: region in space where there is a

maximum probability of finding an electron

Each orbital can hold 2e-s in opposite directions


PAGE 4 OF 30

When e-s are placed in a set of orbital of equal energy,

they occupy them singly and then pairing takes place

e-s placed in opposite direction: both -vely charge & if

placed in same direction, they’d repel. In opposite

direction they create a spin to reduce repulsion

Completely filled or half filled (i.e. one e- in each orbital)

are more stable (reduced repulsion)

2.5 Subshells

Orbitals 1 3 5 7

Max e-s 2 6 10 14

Aufbau’s principle: method of

showing how atomic orbitals are

filled in a definite order to give

lowest energy arrangement possible

Energy difference between 4 & 3 very small ∴ an e- from 4 can be promoted to half-fill or full-fill 3 orbital, to make atom more stable

When filling, fill 4s before 3d and when removing, also

remove first from 4s

2.6 Shapes of Subshells

-Subshell -Subshell

Spherical shape

Increases in size as P.Q

no. increases

Dumbbell shape

2.7 Ionization Energies (I.E)

1st I.E: energy needed to remove 1 mole of e-s from 1

mole of gaseous atom to form 1 mole of unipositive ions

Each successive I.E is

higher than previous one

because as e-s are

removed, protons > e-s ∴ attraction between

protons and remaining

electrons increases

Successive I.Es have

large jump in their value

when e-s removed from

lower energy shell

Deduce group no. by checking when 1st big jump occurs

2.8 Factors affecting Ionization Energy

Nu

cle

ar

Ch

arg

e +ve charge due to protons in nucleus

Greater nuclear charge greater ionization

energy

Sh

ield

ing

Eff

ect

Inner shells of e-s repel outermost e-s, thus

shielding them from +ve nucleus. The more e-

shells, the greater is the shielding effect

Greater effect lower I.E because lesser

attractive force between nucleus & outer e-s

Ato

mic

Ra

diu

s

Distance from the centre of the nucleus to

the outermost orbit

Greater radius lower I.E; distance of

outermost e- to nucleus is large ∴ less energy needed to remove e-

Sta

ble

Co

nfi

g. High I.E needed to remove e-s from

completely or half-filled orbitals

2.9 General 1st I.E Trends

First Ionization Energy Trends

Down a Group Across a Period

DECREASES

New shells added

Attraction of nucleus to

valence e-s decreases

Shielding effect

increases

INCREASES

Shell no. remains same

Proton no. increases

Effective nuclear charge

increases

Atomic radius decreases

2.10 Trend in 1st I.E across 3rd Period

I.E of Al lower than Mg: e- removed in Al is from higher

energy 3p orbital which is further away from nucleus

than 3s e- being removed from Mg. Nuclear attraction is

less for 3p than 3s ∴ I.E of Al is lower than Mg I.E of S lower than P: e- being removed in P is in a half

filled, more stable 3p orbital whereas in S, the pairing of

electrons in 3p results in increased repulsion ∴ less energy need to remove an e-

Big jump occurs

between I.E 1 & 2

∴ part of 1st gp


PAGE 5 OF 30

2.11 Ionic Radius

Ionic radius: describes the size of an ion

Positive ion: smaller radius than

original neutral atom because

shell no. decreases, screening

effect decreases but the

attraction of nucleus increases.

Negative ion: larger ionic radius

than neutral atom because e-s

added while nuclear charge remains same

Groups 1 to 3 5 to 7

Ion Positive Negative

No. of shells

Proton no. and effective nuclear charge increases

Ionic radius decreases

Negative ions always larger than positive ions in the

same period as they have one more shell

3. CHEMICAL BONDING

3.1 Ionic (Electrovalent) Bonding

Ionic bond is the electrostatic

attraction between oppositely

charged ions.

Structure: giant ionic lattice,

crystalline solids

Have high melting and boiling points

Coordination number: number of oppositely charged

ions that surround a particular ion in an ionic solid

3.2 Covalent Bonding

Covalent bond is the bond formed by the sharing of

pairs of electrons between two atoms.

Bonding electrons: e-s involved in bond formation

Non-bonding electrons or lone pair: pair of valence e-s

that are not involved in bond formation

Covalent compounds are made of molecules which are

held together by weak intermolecular forces

They have low melting and boiling points

3.3 Coordinate/Dative Bonding

Coordinate bond is a covalent bond where both

electrons in the bond come from the same atom

Conditions:

o An atom should have a lone pair of electrons

o An atom should be in need of a pair of electrons

Donor: the atom that supplies the pair of electrons

Acceptor: the atom that accepts the pair of electrons

Coordinate bond is represented by an “” drawn from

the atom donating to towards the atom accepting

Formation of Ammonium ion ( ):

Formation of dimer ( :

o Above 750oC, exists as vapor & covalent molecule AlCl3

o As vapor cools, exists as dimer Al2Cl6

o Bond angle as AlCl3 = 120o

o Bond angle as Al2Cl6 = 109.5o

3.4 Orbital Overlap

For a covalent bond to form, atomic orbitals containing

unpaired valence electrons must overlap each other

S –

S

Sigma

Σ

S –

P

Sigma

Σ

P –

P

Sigma

Σ

P –

P

Pi

Sigma bond has greater overlap ∴ Pi bond cannot exist without a Sigma bond.


PAGE 6 OF 30

3.5 Shapes of Covalent Molecules

Shape and bond angles of molecules depend on:

o The number of pairs of electrons around central atom

o Whether these pairs are lone pairs or bonded pairs

Valence shell electrons are arranged in pairs to minimize

repulsion between themselves

Order of repulsion strength (VSEPR Theory):

Type Shape Angle Example

2 Pairs of e-s

2 bonded Linear 180O CO2

3 Pairs of e-s

3 bonded Trigonal

Planar 120O

BF3

4 Pairs of e-s

4 bonded Tetrahedral 109.5O

CH4

3 bonded

1 lone Pyramidal 107O

NH3

2 bonded

2 lone Angular 104.5O

H2O

5 Pairs of e-s

5 bonded Trigonal

Bipyramid 90O

PF5

6 Pair of e-s

6 bonded Octahedral 90O

SF6

3.6 Hydrogen Bonding

Strongest type of intermolecular force in covalent bonds

For hydrogen bonding to occur, we need:

o Molecule having a H atom bonded to F, O or N

o Molecule having F, O or N atom with lone pair of e-s

3.7 Electronegativity

Ability of a particular atom, covalently bonded, to

attract the bonded pair of e-s towards itself

Electronegativity depends on:

o Radius of atom inversely ∝ electronegativity o Nuclear attraction directly ∝ electronegativity Electronegativity increases across a period because

atomic radius ↓ and nuclear attraction ↑

Electronegativity decreases down a group because

atomic radius ↑ and nuclear attraction ↓

Dipole moment: slight charges on

atoms in a covalent bond due to

differences in electronegativity

3.8 Bonds

Bond energy: energy needed to break one mole of a

given bond in one mole of gaseous molecules

Bond length: distance between the centers of two

nuclei of two adjacent atoms

Double bonds are shorter than single bonds because

double bonds have a greater negative charge density

between the two atomic nuclei hence greater attraction

Bond length depends on radii of the two bonded atoms;

larger the radius, longer the bond length

Strength of the bond depends on the length of the bond

3.9 Polar and Non-Polar

Polar Covalent Bonds

Bonds with slight ionic character

Bond formed with atoms of different electronegativity

Bonding e-s attracted more towards atom with greater

electronegativity ∴ unequal sharing of electrons ∴ molecule develops slight charges = Polar Molecule

Polar molecules have dipoles; electric charges of equal

magnitude and opposite sign

The greater the difference in electronegativity of the

two bonded atoms, the greater is the ionic character

Non-Polar Covalent Bonds

Bond formed between:

o Identical atoms: the electronegativity of both atoms is

the same so pair of electron shared equally

o Symmetrical polyatomic molecules: dipoles of bond

exert equal & opposite effects hence cancel charge

Non-polar molecules have no overall charge

Lone -Lone Lone - Bonded Bonded - Bonded

Longer bond Weaker bond More reactive

Water Ammonia ..

..

..


PAGE 7 OF 30

3.10 Intermolecular Forces

Intermolecular forces: weak forces present between

two covalent molecules

Induced Dipole (Van Der Waals’ Forces)

Very weak forces present between non-polar molecules

Due to constant motion of e-s, at an instant, a non-polar

molecule develops poles due to distortion of electron

density giving rise to instantaneous dipole, which is able

to induce a dipole in the adjacent molecules

Van der Waals forces increase with:

o increasing number of contact points between

molecules; point where molecules come close together

o increasing number of electrons (+ protons) in molecule

Permanent Dipole-Dipole Forces

Weak forces present between polar molecules

Molecules always attracted to charged rod, whether +ve

or –ve because molecules have +ve and –ve charges

3.11 Metallic Bonding

Strong electrostatic forces of attraction between metal

cations and delocalized mobile electrons

Structure: lattice of +ve ions surrounded by mobile e-s

Strength of metallic bond increases with:

o Increasing positive charge on the ions in the lattice

o Decreasing size of metal ions in the lattice

o Increasing number of mobile e-s per atom

3.12 Summary

4. STATES OF MATTER

4.1 Basic Assumptions of Kinetic Theory

Ideal gas: a gas whose volume varies in proportion to

temperature and in inverse proportion to pressure.

Noble gases such as helium and neon approach ideal

behavior because of their low intermolecular forces.

Ideal Gas Laws:

Gas molecules move rapidly and randomly

Distance between gas molecules is greater than

diameter of molecules ∴ volume is negligible No forces of attraction/repulsion between molecules

All collisions between particles are elastic conserved

Temperature of gas related to average of molecules

Conditions at which gases behave ideally:

o High temperature

o Low pressure

Limitations of Ideal Gas Laws:

Real gases do not obey kinetic theory in two ways:

o There is not zero attraction between molecules

o We cannot ignore volume of molecules themselves

Deviations visible at low temp. and high pressure

Molecules are close to each other

Volume of molecules not negligible relative to container

VDW forces present, pulling molecules to each other

Pressure is lower than expected from ideal gas

Effective volume is less than expected from ideal gas

4.2 General Gas Equations

Quantity Unit Conversion

Pressure Pascal 1 1000 Volume m3 1 = 1000 = 1 10 cm3 Temperature OK OC + 273

Standard Conditions: 101KPa and 273oK

Mole Fraction

Partial Pressure of a Gas

4.3 Liquid State

Particles touching but may have gaps

Have to slide past each other in random motion

Enthalpy of fusion: heat energy required to change 1

mole of solid into a liquid at its melting point

↑ no. of electron

↓ nuclear attraction

Faster e-distortion

Stronger force


PAGE 8 OF 30

Heating a solid (melting):

o Energy transferred makes solid particles vibrate faster

o Forces of attraction weaken & solid changes to liquid

Enthalpy of vaporization: heat energy required to

change 1 mole of liquid into a gas at its boiling point

Heating a liquid (vaporization):

o Energy transferred makes liquid particles move faster

o Forces of attraction weaken

o Highest energy particles escape first

o Liquid starts to evaporate – temp. below b.p.

o Forces weaken further – particles move faster & spread

o Liquid boils – temp. at b.p.

The evaporation of a liquid in a closed container

o Constant evaporation from surface

o Particles continue to break away

from surface but are trapped in

space above the liquid

o As gaseous particles collide,

some of them hit the surface

of the liquid again, and become trapped there

o An equilibrium is set up in which number of particles

leaving surface is balanced by number rejoining it.

liquid water molecules ⇌ vapor water molecules o In this equilibrium, there will be a fixed number of the

gaseous particles in the space above liquid.

Vapor pressure: pressure exerted by a vapor in

equilibrium with a liquid.

Vapor pressure increases as:

4.4 Solid State

Ionic Lattice Metallic Lattice Simple Molecular

Macromolecular Lattice:

Diamond:

o High m.p./b.p. - each carbon forms four

covalent bonds

o Hard - tetrahedral structure

o Doesn’t conduct heat or electricity – no free e-

o Used for cutting as is strongest known substance and

has sharp edges

Graphite:

o Three strong (sp2) covalent bonds

o Fourth e- in p orbital ∴ forms a pi bond, forming a cloud of

delocalised electron above and

below the planes

o Layers kept together by weak Van der Waal’s forces

o High m.p./b.p. - strong covalent bonds throughout

o Soft – forces between layers are weak

o Conducts electricity - has delocalized electrons

Silicon(IV) Oxide:

o Each Si is bonded to 4 oxygen atoms, but

each oxygen is bonded to 2 Si atoms

o Sand is largely SiO2

o Similar properties to diamond

Hydrogen Bonded Lattice:

In ice form, water molecules slow down

and come closer together

Due to polarity, molecules form

hydrogen bonds between lone pairs of

oxygen & charge of hydrogens

Each water molecule has 2 H-bonds

They arrange themselves into an open crystalline,

hexagonal structure

Due to large spaces, ice is less dense than water

Effect of Hydrogen Bonding on Physical Properties:

o Relatively high m.p./b.p.: many strong H-bonds

o High viscosity: hydrogen bonding reduces ability of

water molecules to slide over each other

o High surface tension: hydrogen bonds in water exert a

downward force on surface of liquid

o Ice less dense than water: larger spaces between

molecules in hexagonal structure

Simple Molecular Lattice:

Iodine:

o Dark grey crystalline solid; vaporizes into purple gas

o m.p./b.p. are slightly higher than room temp

o Slightly soluble in water; dissolves in organic solvents

o Diatomic molecule formed due to covalent bond

between individual atoms

o Molecules have

weak Van der

Waals forces of

attraction between

them

Temp. ↑ Ek ↑ Forces weaken More vapor


PAGE 9 OF 30

Fullerenes:

Buckminsterfullerenes(C60) Nanotubes

C atoms in pentagonal and

hexagonal rings

C atoms in hexagonal

rings only

Spherical Cylindrical

C60 molecules held

together by VDWs

Structure is rod like due

to continuing rings

Conducts heat and electricity

Very strong and tough

Insoluble in water

High m.p./b.p.

4.5 Ceramics

Ceramic: an inorganic non-metallic solid prepared by

heating one or a mixture of substance(s) to a high temp.

Most ceramic are giant molecular structures

Properties of ceramics:

o High m.p./b.p. and hard – strong covalent bonds

o Don’t conduct electricity/heat – no mobile ions or e-s

o Chemically unreactive – e-s held in covalent bonds

4.6 Recycling

Finite resource: resource which doesn't get replaced at

the same rate that it is used up.

Examples of finite resources: copper, aluminium, glass

Advantage of Recycling: ○ Saves energy ○ Reduces

environmental issues ○ Conserves ore supplies ○ Less

wastage ○ Cheaper than extracting

5. CHEMICAL ENERGETICS

5.1 Energy Change in Reactions

Exothermic Reactions Endothermic Reactions

Energy given out

Surrounding warmer

Bond making

∆ negative

Energy taken in

Surrounding cooler

Bond breaking

∆ positive

Exothermic Reactions Endothermic Reactions

Metal + Acid or Water

Neutralization

Combustion

Condensation & Freezing

Thermal decomposition

Dissolving salts

Photosynthesis

Boiling & Evaporation

Standard Enthalpy Conditions:

o Temperature: 298 or 25 o Pressure: 101 or 1 o Solution Conc.: 1

5.2 Enthalpy Change Definitions

Standard molar enthalpy change of Combustion

∆ Formation

∆ Solution

∆ Hydration

∆ Atomisation

∆ Neutralization

∆ Enthalpy change when

1 m

ole

of

ele

me

nt

or

com

po

un

d i

s co

mp

lete

ly

com

bu

ste

d

1 m

ole

of

com

po

un

d i

s fo

rme

d f

rom

its

ele

me

nts

1 m

ole

of

a s

olu

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isso

lve

d i

n a

so

lve

nt

to

form

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in

fin

ite

ly d

ilu

te

solu

tio

n

1 m

ole

of

ion

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th

e g

as

ph

ase

are

dis

solv

ed

in

w

ate

r

1 m

ole

of

ga

seo

us

ato

m

form

ed

fro

m i

ts e

lem

en

t

1 m

ole

of

an

d

com

bin

e t

o f

orm

1

mo

le o

f

under standard conditions in their standard states

5.3 Bond Energy

Energy needed to break a specific covalent bond

Also how much energy is released when a bond forms

5.4 Calculating Enthalpy Changes

∆ ∆ When substance dissolved in water use & of water ∆ is change in temp.: add –ve or +ve to show rise/fall


PAGE 10 OF 30

5.5 Hess’s Law

The total enthalpy change in a chemical reaction is

independent of the route by which the chemical

reaction takes place as long as the initial and final

conditions are the same.

Reason to use Hess’s Law:

o Std. conditions hard to maintain (e.g. exo/endo)

o Elements don’t always react directly

5.6 Calculating Enthalpy Change of…

…Reaction from Formation

…Formation from Combustion

…Hydration from Anhydrous Salt

…Reaction from Bond Energies

6. ELECTROCHEMISTRY

6.1 Calculating Oxidation Numbers

Ionic Molecules: group number = valence electrons

Covalent molecules:

o Rules:

Atoms in a diatomic molecule; oxidation number

= 0

Oxygen in a compound; oxidation

number = -2

Oxygen as peroxide; oxidation number = -1

1st group elements & hydrogen; oxidation

number = +1

H with highly reactive metal; oxidation number =

-1

o Following these rules, all other atoms in a covalent

bond must balance out the charge

6.2 Redox Reactions

Reaction where both oxidation and reduction occur

Can be shown with changes in oxidation numbers of

elements from the product side to the reactant side

E.g. 6 → 2 4 o 4 → 2 ⟹ gain of negative charge ∴ reduction o 4 → 4 ⟹ loss of negative charge ∴ oxidation

6.3 Balancing Equations

Equation: → Half ionic: 2 2 → 2 → For every 2 iodines, there will be 1 nitrogen

Thus first put in correct ratio for iodine and nitrogen

then balance hydrogens and oxygens

Balanced: 4 2 → 2 3


PAGE 11 OF 30

6.4 Electrolysis

Electrolysis: decomposition of an electrolyte by an

electric current. Electrical energy is used to bring about

a chemical reaction; endothermic

Electrolyte: an aqueous solution of an ionic substance

or a molten ionic salt that conducts electricity due to

mobile ions

Electrodes:

o Rods which help current enter the electrolyte

o Inert electrodes: do not take part in the reaction e.g.

graphite or platinum. Steel/titanium used in industry.

o Reactive electrodes: take part in the reaction

6.5 Products of Electrolysis

7. EQUILIBRIA

Reversible reaction: a reaction in which products can be

changed back to reactants by reversing the conditions

Dynamic Equilibrium: the state of a reversible reaction

carried out in a closed container where the rates of

forward and backward reactions are equal and constant

7.1 Le Chatelier’s Principle

When a chemical system in dynamic equilibrium is

disturbed (conditions changed) it tends to respond in

such a way so as to oppose the change and a new

equilibrium is set up

7.2 Equilibrium Constants

Equilibrium Constant

Expressed in terms of

concentration

Expressed in terms of

partial pressure

Only liquids and gases Only gases

Large value of / ⇒ equi. towards products side Smaller value of / ⇒ equi. towards reactants side / changes only with changes in temperature

The amount of reactants that disappear will always

appear in the products in the same ratio as present in a

balanced equation

7.3 Manufacture of Ammonia

HABER PROCESS FLOWCHART

Methane and steam Fractional distillation of

liquid air ⇌ 3

Hydrogen Nitrogen

Compressed and Heated

3 ⇌ 2 Catalyst: Iron beds

Temp: 500oC

Pressure: 200atm

Ammonia

Molten Electrolyte

At the cathode Metal

At the anode Non-metal

Aqueous Electrolyte

At the

cathode

Higher in

reactivity then H+

H2 formed

Lower in

reactivity then H+

Metal

formed

At the

anode

Halide

present

ConcentratedHalogen

fomed

DiluteOxygen

formed

Halide not

present

Oxygen

Formed


PAGE 12 OF 30

7.4 Manufacture of Sulphuric Acid

CONTACT PROCESS FLOWCHART

Burning fossil fuels Fractional distillation of

air

Sulphur Oxygen

⟶

Filter sulphur dioxide

2 3 ⇌ 2 Catalyst: V2O5

Temp: 500oC

Pressure: 1atm (slight)

⟶ [Oleum]

⟶ 2

Sulphuric Acid

SO3 not dissolved directly into water because reaction

explosive and causes H2SO4 to vaporize

Forward reaction exothermic ∴ temp. not too high so rate of backward reaction doesn’t increase & not too

low so particles have for collision and catalyst works

Since reaction highly exothermic, gases must be cooled

No impurities otherwise catalyst will be poisoned

Atmospheric pressure enough because equilibrium

already favours product side

7.5 Acid-Base Equilbria

Brønsted-Lowry Theory:

o An acid is a proton (H+) donor

o A bases is a proton (H+) acceptor

Amphoteric: substances that can act like bases or acids

7.6 Conjugate Pairs

When acid-base reacts, an equilibrium mixture is

formed

is a conjugate acid of base & vice versa

is acid-I and is base-I is a conjugate acid of base & vice versa is acid-II and is base-II

7.7 Strong and Weak Acids and Bases

Strong acids/bases: acids/bases which dissociate almost

completely in solutions

Weak acids/bases: acids/bases which are only partially

dissociated in solutions

Strong and weak acids and bases can be distinguished

by the pH value of their aqueous solutions

Monoprotic acids: donate one H+ proton per molecule

Diprotic acids: donate two H+ protons per molecule

8. REACTION KINETICS

Rate of a reaction: change in concentration of reactants

or products per unit time

Activation energy: minimum energy colliding particles

must possess for a successful collision to take place

Catalysis: acceleration of a chemical reaction by catalyst

8.1 Effect of Concentration Changes

Increasing conc. of reactants increases rate of reaction:

more particles per unit volume, collision rate between

reacting particles increases, ∴ rate of successful collision increases, resulting in increased rate of reaction.


PAGE 13 OF 30

8.2 Maxwell-Boltzmann Theory

Explains effect of temp. & catalyst on rate of reaction

Based on distribution of energy among reacting

molecules under different conditions

8.3 Effect of Temperature

Number of collisions and chance of success will increase

8.4 Effect of Catalyst

Catalyst: a substance that increases rate of reaction but

remains chemically unchanged itself at the end

Does not alter the chemical composition of substances

and only lowers the activation energy

It provides a new route or mechanism to follow for

reactants that requires less energy

Curve unchanged, only activation energy changes

Homogeneous catalysts: reactant and catalyst are in

the same physical state

Heterogeneous catalysts: reactant and catalyst are in

different physical states

Enzymes: a protein molecule that is a biological catalyst.

Most are specific to a substrate & function as lock-key

CI

E AS-

LEVE

L CHE

MIST

RY//

9701

PA

GE

14

OF

30

9.1

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S

SO

2(g

)

SO

3(g

) B

urn

s b

lue

fla

me

+

4

+6

S

. a

cid

ic

9.3

Re

act

ion

of

Na

& M

g w

ith

Wa

ter

Na

& W

ate

r 2

Na

(s) +

2H

2O

(l)

2N

aO

H(a

q) +

H2

(g)

Ve

ry f

ast

, fl

oa

ts,

form

s b

all

& d

isso

lve

s

Mg

& W

ate

r M

g(s

) + 2

H2O

(l)

Mg

(OH

) 2(a

q) +

H2

(g)

Ve

ry s

low

Mg

& S

tea

m

Mg

(s) +

H2O

(g)

Mg

O(s

) + H

2(g

) V

ery

fa

st

9.2

Re

act

ion

of

Ox

ide

s w

ith

Wa

ter

Re

act

ion

O

xid

. N

atu

re

Na

2O

(s)

Na

2O

(s) +

H2O

(l)

2N

aO

H(a

q)

+1

S

. A

lka

lin

e

Mg

O(s

) M

gO

(s) +

H2O

(l)

Mg

(OH

) 2(a

q)

+2

W

. A

lka

lin

e

Al 2

O3

(s)

NO

RE

AC

TIO

N

SiO

2(s

)

P2O

3(s

)

P2O

5(s

)

P2O

3(s

) + 3

H2O

(l)

2H

3P

O3

(aq

)

P2O

5(s

) + 3

H2O

(l)

2H

3P

O4

(aq

)

+3

+5

S

. A

cid

ic

SO

2(g

)

SO

3(g

)

SO

2(g

) + H

2O

(l)

H2S

O3

(aq

)

SO

3(g

) + H

2O

(l)

H2S

O4

(aq

)

+4

+6

S

. A

cid

ic

9.4

Aci

d-B

ase

Re

act

ion

s A

lum

iniu

m o

xid

e i

s a

mp

ho

teri

c ∴ r

ea

cts

wit

h a

cid

an

d b

ase

Al 2

O3 +

H2S

O4

Al 2

(SO

4) 3

+ H

2O

|

A

l 2O

3 +

Na

OH

N

aA

lO2 +

H2O

S

ilic

on

dio

xid

e i

s a

cid

ic:

SiO

2 +

Na

OH

(h

ot

& c

on

c.)

Na

2S

iO3

S

ulp

hu

r d

iox

ide

an

d t

rio

xid

e a

re s

tro

ng

ly a

cid

ic

W

ith

P

rod

uce

s

SO

2(g

) N

aO

H

Na

HS

O3

(aq

)

SO

2(g

) E

xce

ss N

aO

H

Na

2S

O3

(aq

) + H

2O

SO

3(g

) N

aO

H

Na

HS

O4

(aq

)

SO

3(g

) E

xce

ss N

aO

H

Na

2S

O4

(aq

) + H

2O

9. C

HE

MIC

AL

PE

RIO

DIC

ITY

G

rou

p

1

2

3

4

5

6

7

0

Ele

me

nt

So

diu

m

Ma

gn

esi

um

A

lum

iniu

m

Sil

ico

n

Ph

osp

ho

rou

s S

ulp

hu

r C

hlo

rin

e

Arg

on

Ch

ara

cte

r M

eta

l M

eta

lloid

N

on

-me

tals

Str

uct

ure

G

ian

t m

eta

llic

la

ttic

e

Ma

cro

mo

lecu

lar

Sim

ple

mo

lecu

lar

cov

ale

nt

Sim

ple

ato

ms

Bo

nd

ing

M

eta

llic

bo

nd

be

twe

en

ca

tio

ns

an

d d

elo

cali

zed

e-

Co

va

len

t b

on

ds

be

twe

en

ato

ms

Intr

a =

co

va

len

t

Inte

r =

we

ak

VD

Ws

Ato

ms

he

ld b

y

VD

Ws

Dia

gra

m


PAGE 15 OF 30

9.5 Reactions of Elements with Chlorine

Formula Structure Oxid. Nature

Na NaCl(s) Giant ionic

+1 Neutral

Mg MgCl2(s) +2 Neutral

Al AlCl3(s) +3 Acidic

Si SiCl4(l) Simple

molecular

+4 S. Acidic

P PCl3(l)

PCl5(l)

+3

+5 S. Acidic

9.6 Reactions of Chloride with Water

Reaction Nature

NaCl(s) NaCl(s) + H2O(l) NaCl(aq) Neutral

MgCl2(s) MgCl2(s) + H2O(l) MgCl2(aq) W. Acidic

AlCl3(s) AlCl3(s) + H2O(l) Al2O3(s) + HCl(g) Acidic

SiCl4(l) SiCl4(l) + H2O(l) SiO2(s) + HCl(g) S. Acidic

PCl3(l)

PCl5(l)

PCl3(l) + H2O(l) H3PO3(aq) + HCl(g)

PCl5(l) + H2O(l) H3PO4(aq) + HCl(g) S. Acidic

Sodium chloride simply dissolves in water. Water is polar

∴ positive Na+ attracted to OH- while Cl- attracted to H+ MgCl2 slightly acidic because Mg ion has smaller radius &

higher charge ∴ attraction to water is so strong that H2O loses a proton and solution becomes slightly acidic

9.7 Atomic Radius

P+ in nucleus increases so nuclear charge increases

There are more e-, but increase in shielding is negligible

because each extra e- enters same principal energy level

∴ force of attraction between nucleus & e- increases ... So atomic radius decreases.

9.8 Ionic Radius

Ionic radius decreases across a period however, since

non-metals gain electrons, they have one more shell

than metals therefore they always have a larger radius

than metal ions

9.9 Melting Point

Na Al m.p. increases because delocalized e- per atom

increases making metallic bond stronger

Si has highest m.p. due to giant covalent structure

The larger the molecule size, the stronger the VDW

forces ∴ S8 > P4 > Cl2 > Ar 9.10 Electrical Conductivity

Na < Mg < Al because no. of delocalized electrons which

can carry charge increases

Silicon is a semi-conductor

Non-metals – covalent ∴ no charge

AlCl3

anhydrous -covalent; exists as dimer Al2Cl6

Hydrated - exists as ions; Al3+ and

Cl-

AlCl3(s) + H2O(l) Al2O3(s) + HCl(g)

[Al(OH)(H2O)5]2+ + HCl

Oxi

de

+ H

Cl (g

)


PAGE 16 OF 30

9.11 Electronegativity

Increases across period because the bonded e- are in

the same energy level but are attracted more strongly

as no. of protons increases

9.12 First Ionization Energy

Generally increases as no. of protons increases

Decrease Mg Al: more distant and less effective

nuclear charge on 3p orbital

Decrease P S: in S, one electron paired ∴ causing repulsion and easier to lose electron

10. GROUP II – ALKALINE EARTH METALS

m.p./b.p. decreases down group: atoms/ions get larger,

distance between nuclei & e-s increases ∴ bonds weaker m.p./b.p. higher in gp. 2 than 1: 2e-s per atom donated

into delocalized system ∴ metallic bonding stronger density increases down group: mass of atoms increases

faster than their size (volume) as atomic no. increases

10.1 Reactivity of Alkaline Earth Metals

Atomic radius increases down group

Ionisation energy (I.E) decreases down the group

The lower the I.E, easier to remove electrons

Hence metals more reactive down the group

Gp. 2 less reactive than gp. 1 since they need to lose

two e-s ∴ total I.E = 1st I.E. and 2nd I.E. Gp. 2 metals form ionic compounds

10.2 Reaction of Gp. 2 Metals with Oxygen

M s O2 g →2MO s All gp. 2 metals tarnish in air forming oxide coatings

Burn vigorously in oxygen forming white solids

10.3 Reactions with Water

Metals: M s H2O l →M OH 2 aq H2 g Metal Oxides: MO s H2O l →M OH 2 aq Solubility of M, MO and M(OH)2 increases down group

Alkalinity of solution increases down the group

Solubility of M and MO increases do+wn the group

Solubility of M(OH)2 and MSO4 decreases down group

10.4 Reaction with Acid

M s Acid aq → Salt Hydrogen MO s Acid aq → Salt Water

M OH x s Acid aq → Salt Water MCO3 s Acid aq → Salt Water Carbon Dioxide

10.5 Thermal Decomposition of Gp. 2 Metals

MCO3 s∆H MO s CO2 g

2M NO3 2 s∆H 2MO s 2NO2 g O2 g

NO2: thick brown, acidic and soluble gas

Thermal stability increases down the group ∴ decomposition becomes more difficult.

10.6 Uses of Group II Metals

Calcium compounds:

Calcium oxide (lime): basic oxide used to neutralize

acidic soil and used as a drying agent for drying

ammonia

Calcium carbonate (limestone): used as building

material (cement, concrete) etc., for extraction of iron,

glass industry, neutralize soil or chemical waste


PAGE 17 OF 30

11. GROUP 17 – HALOGENS

11.1 Trends in Colour and Volatility

Fluorine Yellow Gas

m.p

. &

b.p

.

incr

ea

ses

Vo

lati

lity

de

cre

ase

s

Chlorine Yellow-Green

Bromine Orange-Brown Liquid

Iodine Grey-Blue Solid

Astatine Black

As atomic number increases, the number of electrons

increases, this increases VDW forces so stronger bonds

thus m.p./b.p. increases and volatility decreases

11.2 Oxidising Ability

Halogens have high electron affinity (they gain electrons

easily) hence they are good oxidising agents

Oxidising ability decreases down the group because

electron affinity decreases as atomic size increases.

11.3 Some reactions of the halide ions

X2(g) + H2(g) 2HX(g)

Product Reaction Description

HF Reacts explosively in all conditions

HCl Reacts explosively in sunlight

HBr Reacts slowly on heating

HI Forms an equilibrium mixture on heating

Thermal stability of halogen hydrides decreases down

the group because:

o Size of halogen atom increases

o ∴ nuclear attraction decreases o The H – X bond becomes longer and weaker

o Thus less energy needed to break the bond

Bond energies decrease down the group

(Sub) Halide ions and aq. Silver Ions

Ag+(aq) + X-(aq) AgX(s)

Halide

Ion

With Silver

Nitrate

With dilute

aq. ammonia

With conc.

aq. ammonia

Cl- White ppt. ppt. dissolves

Br- Cream ppt. ppt. dissolves

I- Yellow ppt.

The solubility of these ppts. are tested with dilute and

conc. aq. ammonia to confirm presence of ion.

If ppt. dissolves, it forms a complex ion:

AgX(s) + 2NH3(aq) [Ag(NH3)2]+(aq) + X-

The complex ion formed is called Diamine Silver(I) ion

∶⟶ ←∶

(Sub) Halide ions and aq. Sulphuric Acid

Metal Halide + Conc. H2SO4(aq) Hydrogen Halide

Conc. H2SO4(aq) is an oxidising agent

This reaction is used for preparation of hydrogen halides

Chlorine NaCl(s) + H2SO4(aq) HCl(g) + NaHSO4(aq)

Bromine NaBr(s) + H2SO4(aq) HBr(g) + NaHSO4(aq)

HBr(g) + H2SO4(aq) Br2(g) + SO2(g) + H2O(l)

Iodine

NaI(s) + H2SO4(aq) HI(g) + NaHSO4(aq)

HI(g) + H2SO4(aq) I2(g) + SO2(g) + H2O(l)

HI(g) + H2SO4(aq) I2(g) + H2S(g) + H2O(l)

11.4 The reactions of chlorine with aqueous

sodium hydroxide

Disproportionation: a reaction in which the same

substance is oxidized and reduced simultaneously

producing two different products

When chlorine reacts with a solution of cold aqueous

sodium hydroxide, the disproportionation goes to lower

oxidation states

Cl 2NaOH → NaCl NaClO H O With a hot solution, the oxidation state of chlorine goes

up to V 3Cl 6NaOH → NaClO 5NaCl 3H O

This happens as the chlorate is formed by

disproportionation of hypochlorite and hypochlorous

acid

ClO 2HClO → ClO 2HCl Higher temperatures promotes the formation of

hypochlorous acid through hydrolysis of hypochlorite

and therefore speeds up the reaction

11.5 Some important uses of halogens and of

halogen compounds

Fluorine:

o To make chlorofluorocarbon (CFCs)

o As fluoride in toothpaste

o To make polytetrafluoroethylene (PTFE) – non sticking

coating in pots and pans

Bromine and Iodine: manufacture of photographic films

Chlorine:

o In bleaches

o To make PVC and chlorofluorocarbon (CFCs)

o As solvents


PAGE 18 OF 30

Use of chlorine in water purification:

o The oxidising power of chlorine is used in treatment of

water to kill bacteria

Cl2(aq) + H2O(l) HCl(aq) + HClO(aq)

HClO(aq) HCl(aq) + O

o This disproportionation reaction produces reactive

oxygen atoms which kill bacteria

12. NITROGEN AND SULPHUR

12.1 Lack of Reactivity of Nitrogen

Nitrogen molecule has three strong covalent bonds

Bond is very strong and requires high energy for

splitting the two nitrogen atoms of a molecule.

It reacts only under extreme temperature or pressure or

in presence of catalyst.

12.2 Ammonium

Lone pair of e-s of nitrogen forms a coordinate bond

with the H+ ion

Formation: NH3(g) + H+ NH4+

Shape: tetrahedral

Bond angle: 109.5o

Bond length: equal lengths

Displacement of ammonia from its salts:

Any

Ammonium

Salt

+ Any

Base

warm Ammonia

Gas + Salt + Water →

12.3 Uses of Ammonia & its Compounds

Used in the production of nitric acid

Used in the production of inorganic fertilizers

Used in the production of nylon

Used in the production of explosives

12.4 Eutrophication

Nitrate fertilisers leach into rivers and lakes after rain

Water plants grow more than usual

They block sunlight and kill plants underneath

Bacteria/fungi decompose remains using the O2

Fish and other creatures die from oxygen starvation

12.5 Oxides of Nitrogen

N2(g) + O2(g) 2NO(g) or N2(g) + ½O2(g) NO2(g)

Naturally: during lightning, EA provided for N2 to react

Man-made: in car engine, high temp. and pressure

Catalytic convertors: exhaust gases passed through

catalytic convertors containing a catalyst (platinum/

palladium/nickel) helping to reduce oxides to nitrogen.

Catalytic role in oxidation of sulphur dioxide:

12.6 Pollution

Acid Rain: SO3 + H2O H2SO4

2NO2 + H2O HNO3 + HNO2 or NO2 + H2O + ½O2 HNO3

Damages trees & plants, kills fish and other river life,

buildings, statues and metal structures

Combustion Pollutants:

Nitrogen oxide (NO): formed by reaction of N2 and O2 in

the engine, forms acid rain and respiratory problems

Carbon monoxide (CO): source: incomplete combustion

of hydrocarbon fuel, toxic effect on haemoglobin

12.7 Food Preservation

SO2 is used by itself or as a sulphite to preserve food

SO2 + H2O H2SO3(aq)

SO2 & sulphites inhibit growth of bacteria, yeasts, etc. &

are reducing agents, so reduce rate of oxidation of food.

Used to prevent spoilage of dried fruit, dehydrated

vegetables and fruit juices.

13. INTRODUCTION TO ORGANIC CHEMISTRY

Organic chemistry: study of hydrocarbons and their

derivatives

Carbon can form a variety of compounds because:

o Carbon is tetravalent

o Carbon-carbon bonds can be single, double or triple

o Atoms can be arranged in chains, branches and rings

Homologous series: a series of compounds of similar

structures in which:

o contain the same functional group

o all share same general formula

o formula of homologue differs from neighbour by CH2

o similar chemical properties

o gradual change in physical properties as Mr increases

Functional group: an atom or group of atoms in an

organic molecule that determine the characteristic

reactions of a homologous series.

Alkyl group: a reactive group which is alkane minus 1 H


PAGE 19 OF 30

13.1 Hybridization

Hybridisation: mixing up of different atomic orbitals

resulting in new orbitals of equal energy.

Carbon’s electron configuration:

Ground State Excited State

2s 2p 2s 2p

↿⇂ ↑ ↑ ↑ ↑ ↑ ↑

sp3 sp2 sp

All orbitals mix 2s, 2px, 2py mix 2s and 2px mix

4 sp3 orbitals 3 sp2 orbitals

1 pure p orbital

2 sp orbitals

2 pure p orbitals

Ratio of characteristics s : p

1 : 3 1 : 2 1 : 1


PAGE 20 OF 30

13.2 Classes of Compound

Organic Family Suffix Example

Alkanes

-ane

Methane

Alkenes

-ene

Ethene

Halogenoalkanes

halo- … -ane

Chloroethane

Alcohols

-ol

Methanol

Aldehydes

-al

Methanal

Ketones

-one

Propanone

Carboxylic Acid

-oic

Methanoic acid

Esters

-oate

Methyl ethanoate

Amines

-amine

Methylamine

Nitriles

-nitrile

Ethyl nitrile


PAGE 21 OF 30

13.3 Types of Formulae

Hexane

Displayed Formula Structural Formula

CH3-CH2CH2CH2CH2-CH3

or

CH3(CH2)4CH3

Skeletal Formula Molecular Formula

C6H14

13.4 Nomenclature

Select longest chain as the main chain

Other carbon chains as substituent alkyl groups

Give lowest number C in main chain to substituent

If different alkyl groups present on identical position,

give simpler alkyl smaller number

Two or more alkyl groups present, order alphabetically

If same substituent repeated use di, tri, tetra prefix

If ring of carbon present, use prefix cyclo

Write position of double bond in alkene e.g. but-1-ene

13.5 Breaking of Covalent Bonds

Homolytic Fission:

Two atoms sharing e- pair of similar electro-tivity

When bond breaks, each atom takes one e- from pair of

electrons forming free radicals

Free radicals: electrically neutral atoms or group of

atoms with unpaired electrons very reactive

Free radical reaction catalysed by heat or light

Heterolytic Fission:

Two atoms sharing e- pair are of different electro-tivity

When bond breaks, one of the bonded atoms takes both

bonding e-s

Results in formation of +ve and –ve ions

If +ve charge on C, its called carbocation or carbonium

If –ve charge on C, its called carbanion

Note: homolytic fission require less energy than heterolytic

13.6 Types of Reagents

Nucleophilic reagent (nucleophile): donator of pair of e-

Must have lone pair of e-s

Attack centre of +ve charge (positive pole)

Reaction with nucleophile called nucleophilic reactions

Examples: CH-, Cl-, NH3, H2O, CN-

Electrophilic reagent (electrophile): acceptor of pair of e-

+ve ions or e- deficient molecules

Attack regions of high e- density

Examples: Br+, CH3+, AlCl3

13.7 Types of Reaction

Addition reaction: single product formed

o Electrophilic addition (alkenes)

o Nucleophilic addition (carbonyl compounds)

Substitution reaction: two products formed

o Nucleophilic substitution (halogenoalkanes)

o Free radical substitution (alkanes)

Elimination reaction: more than one product formed,

small molecule removed from reactant (alcohols and

halogenoalkanes)

Hydrolysis reaction: breaking down of molecule by

water, sped up by acid or alkali (esters and alkenes)

13.8 Oxidation and Reduction

Oxidation: addition of oxygen or removal of hydrogen

Reduction: addition of hydrogen or removal of oxygen

13.9 Shapes of Ethane and Ethene

Ethane Ethene

sp3 bonds

all sigma bonds

Planar Shape

H – C – H bond = 120o

Benzene

13.10 Isomerism

Existence of two or more compounds with the same

molecular formula but different structural formula

ISOMERS

Structural: molecules that have the same molecular formula, but have a different arrangement of the atoms

Chain Position Functional

Stereo: molecules that have same structure, but different spatial

arrangement of atoms in 3D space

Geometric Optical


PAGE 22 OF 30

Note:

Straight chain alkanes have higher b.p. than branched

Branching makes molecule more spherical reduces

contact points VDW forces decreases

13.11 Chain Isomers

Isomers have different carbon chain length

Same chemical properties but slightly different physical

Example:

13.12 Position Isomers

Isomers differ in position of substituent atoms or group

or the functional group

Same chemical properties but slightly different physical

Example:

But-1-ene

But-2-ene

13.13 Functional Isomers

Isomers have different functional groups, belong to

different homologous series

Have different physical and chemical properties

Ratio of C : H Functional Gps. Example

1 : 3 Alcohol & Ether C2H6O

1 : 2 Aldehyde & Ketone C3H6O

1 : 2

Must have O2

Carboxylic acid &

Ester C3H6O2

13.14 Geometric (cis/trans) Isomers

Shown only by alkenes

Arises due to restriction of double bond

Only possible when each carbon has 2 different groups

cis-trans isomers have different b.p.

cis isomers have higher dipole

trans isomer of symmetrical alkene has zero dipole

13.15 Optical Isomers

Arises from different arrangement of atoms or groups in

3D space resulting in two isomers

Have effect on polarised light

Chiral carbon: a carbon having 4 single bonds and 4

different atoms or groups

Isomers non-super-imposable images of each other

Have same physical and chemical properties

No. of optical isomers in a molecule containing chiral

carbons 2

14. HYDROCARBONS

14.1 Properties

Generally unreactive:

All C–C bonds single; alkanes = saturated hydrocarbons

Non-polar ∴ no center of charge to act as either nucleophile or electrophile ∴ cannot attract polar reagents like acids, bases, metals or oxidizing agents

Physical properties:

The volatility of the alkanes decreases and m.p/b.p

increases as number of carbon atoms increases

Reason: increasing Van der Waals forces

14.2 Combustion

Used as fuel because they burn in oxygen to given out

large amounts of energy

Alkanes kinetically stable in presence of O2; combustion

occurs when necessary amount of Ea supplied

Reaction occurs only in gas phase

Complete: carbon dioxide + water

Incomplete: carbon monoxide + carbon (soot) + water

General Equation of Hydrocarbon Combustion:

4 → 2


PAGE 23 OF 30

14.3 Substitution

Alkanes react with halogens: and

Example: Chlorination of Methane

Reagent Condition Reaction

Type Mechanism

Cl2(g) UV light Substitution Free Radical

Initiation:

o Energy of a photon of light absorbed

o bond breaks homolytically Propagation:

o Highly reactive ∙ collides with a molecule forming a new free radical; ∙

⋅→ ⋅ ⋅ → ⋅ o This can then react with another and process

repeats if sufficient present until all are replaced

Termination:

o Reaction ends when 2 free radicals collide & combine

⋅ ⋅→ ⋅ ⋅→ Products: forms large amounts of CH3Cl and HCl and

small amount C2H6; separated by fractional distillation

Products and free radicals differ due to:

o Halogen used: bromine requires more light

o Alkane used: ↑ no. of C = ↑ variety of products

14.4 Cracking

Breaking of large less useful alkanes into useful, more

energy value smaller products using heat & catalyst

Products: smaller alkanes and alkenes or

smaller alkenes and hydrogen gas

Thermal cracking: high temp. & pressure

Catalytic cracking: high temp. & catalyst

14.5 Hydrocarbons as Fuels

Source of alkanes: crude oil

Steady change in b.p. of alkanes allows crude oil to be

separated by fractional distillation

Catalytic conversion of CO and NOx:

o 2NO2 + 4CO N2 + 4CO2

o 2NO + 2CO N2 + 2CO2

14.6 Alkenes

Unsaturated hydrocarbons

Contain at least one C=C double bond

General formula: CnH2n (like cycloalkanes)

Source of alkenes:

o Cracking alkanes

o Dehydration of alcohols

More reactive than alkanes due to presence of double

bond; pi electrons loosely and more susceptible to

attacks by e- deficient groups like electrophiles

Alkenes combust completely carbon dioxide + water

Give energy but not used as fuels; have other uses

14.7 Electrophilic Addition Mechanism

Electrophile forms by heterolytic fission

Electrophile attacks double bond

Pair of e-s from double bond migrate to electrophile and

bond breaks

Carbocation formed which attacks the nucleophile

14.8 Carbocations

Markovnikov’s principle: an electrophile adds to an

unsymmetrical alkene so that the most stable

carbocation is formed as an intermediate

Hydrogen binds to carbon with more hydrogens

Inductive effect of alkyl groups:

o Alkyl groups donate e- to the ring

o Producing a positive inductive effect

o A larger alkyl group has a weaker inductive effect

Hy

dro

-

ha

log

en

ati

on


PAGE 24 OF 30

14.9 Addition Reactions H

yd

rog

en

ati

on

Alkene + H2 Alkane

Reagent: H2(g)

Condition:

o Catalyst: Nickel

o Temp.: 100oC

o Press.: 2 atm.

Use: convert liquid oils to saturated solid fats

Alkene + X2 Dihaloalkane

Reagent: Halogen(aq)

Condition: r.t.p./dark

Alkene + Hydrohalogen Halogenoalkane

Reagent: Hydrohalogen(g)

Condition: r.t.p.

Alkene + H2O(g) Alcohol

Reagent: steam

Condition:

o Catalyst: H3PO4 – phosphoric acid

o Temp.: 300oC

o Press.: 70atm

14.10 Oxidation of Alkenes

Both oxidation and addition to double bond involved

KMnO4 changes from pink to colourless

With Cold Dil. Acidified KMnO4/H+

Diol is formed

With Hot Conc. Acidified KMnO4/H+

Leads to the rupture of the double bond

Two compounds are formed

Products formed depend on alkene

14.11 Polymerization

Repeated addition of 1000s of alkene molecules

(monomer) to each other forming a macromolecule

Polyethene:

o LDPE: cling wrap

o HDPE: water pipes, wire

insulation

Polychloroethene (PVC):

o Water pipes

o Insulation of wires

General conditions: high

pressure,

high temperature and catalyst

Disadvantages:

o Non-biodegradable: does not break down so increases

amount of volume needed for landfill sites

o Combustion produces harmful gases which contribute

to global warming e.g. SO2, CO2 and HCl from PVCs

Disposal of Polymers:

o Recycle existing plastic

o Make polymers biodegradable by adding starch units

Ha

log

en

ati

on

H

yd

rati

on


PAGE 25 OF 30

15. HALOGEN DERIVATIVES

15.1 Types of Halogenoalkanes

Primary 1o Secondary 2o Tertiary 3o

SN2 SN1

15.2 Strength of C – Hal Bond

Polar Nature Bond Energy Reactivity

Fluoro

De

cre

ase

De

cre

ase

Incr

ea

ses

Chloro

Bromo

Iodo

Electro-tivity

decreases

down group

Bond length increases, bond

energy decreases, lower EA so

more reactive

15.3 Nucleophilic Substitution Mechanism

The C – X bond is a polar bond, has partial charges due

to high electro-tivity of halogen.

The carbocation is easily susceptible to attack by a

nucleophile

SN1 Mechanism:

Unimolecular – only one molecule involved in 1st step

Tertiary halogenoalkanes

SN2 Mechanism:

Bimolecular – two molecules involved in 1st step

Primary and secondary halogenoalkanes

15.4 Nucleophilic Substitution Reaction

Hy

dro

lysi

s

R – X + OH- R – OH + X-

Reagent: strong alkali; NaOH(aq) or KOH(aq)

Condition: heat/reflux

Fluoroalkanes are not hydrolysed because the

C – F bond is too strong

Ease of hydrolysis increases:

Primary < Secondary < Tertiary

Tertiary halogenoalkanes can be hydrolysed

without alkali

Note: if any Cl- or Br- ions present in NaOH(aq),

these ions will interfere with reaction

Nit

rile

(cy

an

ide

)

R – X + CN- RCN + X-

Reagent: KCN or NaCN in ethanol

Condition:

o Solvent: Ethanol

o Heat/Reflux

Reaction forms a C – C bond therefore no. of

C increases; name has one more carbon

Pri

ma

ry A

min

es R – X + NH3 RNH2(l) + HX(g)

Reagent: Ammonia (NH3)

Condition: ammonia in alcohol under

pressure in sealed container

Note: if excess conc. ammonia used, HX

reacts with it forming NH4X

15.5 Reflux

Many organic reactions proceed slowly

Heating done under reflux to prevent

volatile organic solvents to evaporate

Mechanism similar to simple distillation

15.6 Elimination Reaction

R – X + OH- Alkene + X- + H2O

Mechanism:

Reagent: ethanolic NaOH or KOH

Conditions: temp. 60oC, reflux

OH- acts as a proton acceptor; it accepts the H+ loss from

the halogenoalkanes during elimination

Elimination become progressively more easier

Primary < Secondary < Tertiary

Note: the carbon atom adjacent to carbon with halide

must have at least one hydrogen attached to it.


PAGE 26 OF 30

15.7 Uses of Halogenoalkanes

CFCs are inert and can be liquefied easily: Strength of

C – X bond is very high, hence do not decompose easily

and are not flammable.

Uses:

o As propellants in aerosol cans

o As solvents in dry-cleaning

o As refrigerant for freezers and fridges

o Fire extinguishers, insecticides and pesticides

15.8 CFCs Effect on Ozone Layer

Causes the destruction of the ozone layer

CFCs escape in atmosphere and because of their

inertness, remain without further reaction until they

reach the stratosphere and ozone layer.

In stratosphere, high energy U.V causes Cl atom to split

of CFC molecule forming Cl⋅ which reacts with ozone This is a catalytic cycle where one Cl⋅ can react with

many O3 thus causing destruction of ozone layer:

Cl⋅ + O3(g) ⋅OCl(g) + O2(g) ⋅OCl(g) + O(g) Cl⋅ + O2(g)

Can react and breakdown another O3 molecule

Note: alternative is using HCFCs (replace Cl with H or

more F atoms) as they break down more easily and do not

release Cl less effect on ozone layer

16. HYDROXY COMPOUNDS

16.1 Types of Alcohols

Primary 1o Secondary 2o Tertiary 3o

Source of Alcohols:

o Hydration of alkenes

o Fermentation

16.2 Properties

Physical Properties:

Colourless liquids at r.t.p

b.p. and density increases with increasing C atoms and

also with increasing OH groups

Boiling Point:

b.p. decreases

because:

b.p. of alcohols > alkenes as they have hydrogen bonds

Solubility of Alcohols in Water:

Smaller alcohols mix completely with water since strong

hydrogen bonds occur between alcohols and water

As hydrocarbon nature increase (i.e. more C-C… bonds),

the non-polar character outweighs the ability of the OH

to form hydrogen bonds and ∴ solubility decreases Small alcohols (e.g. ethanol) are good solvents for both

polar and non-polar compounds as they have polar and

non-polar components

16.3 Reaction with Sodium

R – OH + Na(l) RO- Na+ + ½ H2(g)

Type of reaction: acid-base

Reagent used: liquid sodium metal

Reactivity of alcohols decreases with increasing chain

lengths of hydrocarbon

Reaction less vigorous than that of Na and water which

shows water is a stronger acid than alcohol

16.4 Reaction with Carboxylic Acids

Reagent Condition Type of Reaction

R-COOH Heat-reflux

Conc. H2SO4 Esterification

Primary Secondary Tertiary

branching increases VDWs decrease b.p. decreases


PAGE 27 OF 30

Naming esters:

Properties of Esters:

Esters are volatile compounds – no H-bonds so low m.p.

Polar molecules – soluble in organic solvents

Sweet, fruity smelling liquids

Many occur naturally e.g. as fats, oils & flavours in fruits

Used in food flavourings and perfumes and as solvents

16.5 Hydrolysis of Esters

16.6 Dehydration of Alcohols

Alcohol(l) Alkene + H2O(l)

Condition Type of Reaction

Conc. H2SO4

H3PO4 at 180oC

Al2O3 at 300oC

or

Elimination or

Mechanism:

Adjacent carbon to carbon with OH must have at least

one hydrogen (tertiary cannot undergo dehydration)


PAGE 28 OF 30

16.7 Halogenation

Type of Reaction: Nucleophilic Substitution

R – OH R – X

Forming Reagent Producing:

R –

OH

+

Reactions Condition

Alkyl Chlorides

Conc. HCl RCl(l) + H2O Zn + Heat/Reflux

SOCl2 RCl(l) + SO2(g) + HCl(g) r.t.p

PCl5 RCl(l) + POCl3(aq) + HCl(g)

PCl3 RCl(l) + H3PO3(aq) + HCl(g) Heat/Reflux

NaBr + H2SO4(aq) HBr Alkyl Bromides

HBr(g) RBr(l) + H2O r.t.p

P + Br2 –warm PBr3 PBr3(g) RBr(l) + H3PO3(aq)

P + I2 –warm PI3 Alkyl Iodide PI3(g) RI(l) + H3PO3(aq) r.t.p

16.8 Oxidation of Alcohols

Reagent: Oxidising agents

Reagent Type of Reaction

Acidified K2Cr2O4

Orange to Green

Acidified KMnO4

Pink to Colourless Oxidation

Primary Alcohol

R-CH2-OH

Secondary Alcohol

R-CH(OH)-R

If aldehyde needed, must distil

as soon as it forms

Tertiary alcohols not oxidised because no hydrogens

attached to carbon with OH group so oxidising agent

colour does not change

16.9 Tests for Alcohols

Reagent Result with:

Primary Secondary Tertiary

Na metal Bubble of H2 Gas

K2Cr2O4/H+ Green

KMnO4/H+ Colourless

17. CARBONYL COMPOUNDS

Boiling Point:

Solubility:

Smaller carbonyl compounds: completely soluble as

they form hydrogen bonds with water molecules; are

good solvents for polar & non-polar solutes

Larger carbonyl compounds: polar nature decreases

and non-polar nature increases; ability to form

hydrogen bonds decreases

17.1 Nucleophilic Addition with HCN

Reagent Condition Type of Reaction

HCN HCN w/alkali or

HCN w/KCN Nucleophilic Addition

Since HCN added, carbon chain increases

Product formed is hydroxynitrile or cyanohydrine

Aldehydes are more susceptible to nucleophilic attacks

than ketones

Smaller carbonyl compounds more reactive

Product has a chiral carbon ∴ exhibits optical isomerism Mechanism:


PAGE 29 OF 30

Note: HCN is a poor nucleophile and with few CN- ions,

the reaction is slow. To increase CN- conc.:

Make HCN react in presence of alkali

HCN + OH- H2O + CN-

Addition of KCN and dilute H2SO4 can provide HCN and

more CN- ions

17.2 Reduction of Carbonyl Compounds

Type of Reaction: nucleophilic addition (H- ions)

Reducing agents:

o NaBH4 – sodium tetrahydrioborate

o LiAlH4 – lithium aluminium hydride

o H2/Pt or Ni

Aldehydes ⟹ 1o Alcohols Ketones ⟹ 2o Alcohols R-CHO + 2[H] RCH2OH R-CO-R + 2[H] R-CH(OH)-R

17.3 Testing Carbonyl Compounds

2,4,- dinitrophenylhydrazine:

It is a nucleophilic addition & condensation/elimination

Forms: red/orange ppt.

The m.p. of the ppt. can be used to identify individual

aldehydes and ketones

Tests Given only by Aldehydes:

Tollen’s Reagent

Solution of AgNO3 + aq. NH3 excess [Ag(NH3)2]+

Aldehyde + Tollen’s Reagnet Silver Mirror

Ag+ reduced to Silver and –CHO oxidised to acid

2Ag+ + RCHO 2Ag + RCOOH- + H+

Fehling’s Solution

CuSO4 in ammonia solution

Aldehyde + Fehling’s Solution Red ppt.

Cu2+ reduced to Cu(I) oxide and –CHO oxidised to acid

2Cu2+ + RCHO 2Cu+ + RCOOH- + H+

18. CARBOXYLIC ACIDS AND DERIVATIVES

Weak acids; don’t dissociate completely

Forms hydrogen bonds:

o High m.p./b.p.

o High solubility of smaller carboxylic acids

Forms hydrogen bonded dimers when

pure vapour, liquid or solid & when

dissolved in non-polar organic solvents

18.1 Formation of Carboxylic Acids

From alcohols: complete oxidation of primary alcohols

From aldehydes: oxidation of aldehydes

From nitriles: acid/base hydrolysis of a nitrile

18.2 Formation of Salts

Heterolytic fission of the hydroxyl bond (-OH)

Salts called carboxylates

19. ANALYTICAL TECHNIQUES

19.1 Infra-red Spectroscopy

This is when a sample being analysed is irradiated with

electromagnetic waves in the infra-red region of the

electromagnetic spectrum.

Machine used is spectrophotometer and it detects

intensity of wavelengths of infra-red that pass through

the sample

The energy absorbed corresponds to changes in

vibration of bonds leading to the bond being to stretch,

bend and twist


PAGE 30 OF 30

At a specific frequency, the resonance frequency, the

largest vibrations are obtained

Each type of vibration will absorb characteristic

wavelengths of infra-red radiation

We can hence identify the presence (or absence) of

different functional groups from the absorbance pattern

on an infra-red spectrum

19.2 Monitoring Air Pollution

IR spectroscopy identifies particular bonds in a

molecule, and so each pollutant will show a different

pattern of absorptions – this allows the identification of

the pollution

It is also possible to measure the concentration of each

pollutant with the different amounts of absorption

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